Chapter 28

1. Chapter 28 Atomic Physics

2. Importance of Hydrogen Atom • Hydrogen is the simplest atom • The quantum numbers used to characterize the allowed states of hydrogen can also be used to describe (approximately) the allowed states of more complex atoms • This enables us to understand the periodic table

3. More Reasons the Hydrogen Atom is so Important • The hydrogen atom is an ideal system for performing precise comparisons of theory with experiment • Also for improving our understanding of atomic structure • Much of what we know about the hydrogen atom can be extended to other single-electron ions • For example, He+ and Li2+

4. Sir Joseph John Thomson • “J. J.” Thomson • 1856 - 1940 • Discovered the electron • Did extensive work with cathode ray deflections • 1906 Nobel Prize for discovery of electron

5. Early Models of the Atom 28.1 • J.J. Thomson’s model of the atom • A volume of positive charge • Electrons embedded throughout the volume • A change from Newton’s model of the atom as a tiny, hard, indestructible sphere

6. Early Models of the Atom, 2 • Rutherford, 1911 • Planetary model • Based on results of thin foil experiments • Positive charge is concentrated in the center of the atom, called the nucleus • Electrons orbit the nucleus like planets orbit the sun

7. Rutherford’s Gold Foil Experiment

8. Scattering Experiments • The source was a naturally radioactive material that produced alpha particles • Most of the alpha particles passed though the foil • A few deflected from their original paths • Some even reversed their direction of travel

9. Difficulties with the Rutherford Model • Atoms emit certain discrete characteristic frequencies of electromagnetic radiation • The Rutherford model is unable to explain this phenomena • Rutherford’s electrons are undergoing a centripetal acceleration and so should radiate electromagnetic waves of the same frequency • The radius should steadily decrease as this radiation is given off • The electron should eventually spiral into the nucleus, but it doesn’t

10. Emission Spectra 28.2 • A gas at low pressure has a voltage applied to it • A gas emits light characteristic of the gas • When the emitted light is analyzed with a spectrometer, a series of discrete bright lines is observed • Each line has a different wavelength and color • This series of lines is called an emission spectrum

11. Examples of Emission Spectra

12. Emission Spectrum of Hydrogen – Equation • The wavelengths of hydrogen’s spectral lines can be found from • RH is the Rydberg constant • RH = 1.097 373 2 x 107 m-1 • n is an integer, n = 1, 2, 3, … • The spectral lines correspond to different values of n

13. Spectral Lines of Hydrogen • The Balmer Series has lines whose wavelengths are given by the preceding equation • Examples of spectral lines • n = 3, λ = 656.3 nm • n = 4, λ = 486.1 nm

14. Absorption Spectra • An element can also absorb light at specific wavelengths • An absorption spectrum can be obtained by passing a continuous radiation spectrum through a vapor of the gas • The absorption spectrum consists of a series of dark lines superimposed on the otherwise continuous spectrum • The dark lines of the absorption spectrum coincide with the bright lines of the emission spectrum

15. Applications of Absorption Spectrum • The continuous spectrum emitted by the Sun passes through the cooler gases of the Sun’s atmosphere • The various absorption lines can be used to identify elements in the solar atmosphere • Led to the discovery of helium

16. Absorption Spectrum of Hydrogen

17. Neils Bohr 28.3 • 1885 – 1962 • Participated in the early development of quantum mechanics • Headed Institute in Copenhagen • 1922 Nobel Prize for structure of atoms and radiation from atoms

18. The Bohr Theory of Hydrogen • In 1913 Bohr provided an explanation of atomic spectra that includes some features of the currently accepted theory • His model includes both classical and non-classical ideas • His model included an attempt to explain why the atom was stable

19. Bohr’s Assumptions for Hydrogen • The electron moves in circular orbits around the proton under the influence of the Coulomb force of attraction • The Coulomb force produces the centripetal acceleration

20. Bohr’s Assumptions, cont • Only certain electron orbits are stable • These are the orbits in which the atom does not emit energy in the form of electromagnetic radiation • Therefore, the energy of the atom remains constant and classical mechanics can be used to describe the electron’s motion • Radiation is emitted by the atom when the electron “jumps” from a more energetic initial state to a lower state • The “jump” cannot be treated classically

21. Bohr’s Assumptions, final • The electron’s “jump,” continued • The frequency emitted in the “jump” is related to the change in the atom’s energy (measured as photons) • It is generally not the same as the frequency of the electron’s orbital motion • The frequency is given by Ei – Ef = h ƒ • The size of the allowed electron orbits is determined by a condition imposed on the electron’s orbital angular momentum

22. Mathematics of Bohr’s Assumptions and Results • Electron’s orbital angular momentum • me v r = n ħ where n = 1, 2, 3, … • The total energy of the atom • The energy of the atom can also be expressed as

23. Bohr Radius • The radii of the Bohr orbits are quantized • This is based on the assumption that the electron can only exist in certain allowed orbits determined by the integer n • When n = 1, the orbit has the smallest radius, called the Bohr radius, ao • ao = 0.052 9 nm

24. Radii and Energy of Orbits • A general expression for the radius of any orbit in a hydrogen atom is • rn = n2 ao • The energy of any orbit is • En = - 13.6 eV/ n2

25. Specific Energy Levels • The lowest energy state is called the ground state • This corresponds to n = 1 • Energy is –13.6 eV • The next energy level has an energy of –3.40 eV • The energies can be compiled in an energy level diagram

26. Specific Energy Levels, cont • The ionization energy is the energy needed to completely remove the electron from the atom • The ionization energy for hydrogen is 13.6 eV • The uppermost level corresponds to E = 0 and n  

27. Energy Level Diagram • The value of RH from Bohr’s analysis is in excellent agreement with the experimental value • A more generalized equation can be used to find the wavelengths of any spectral lines

28. Generalized Equation • For the Balmer series, nf = 2 • For the Lyman series, nf = 1 • Whenever an transition occurs between a state, ni to another state, nf (where ni > nf), a photon is emitted • The photon has a frequency f = (Ei – Ef)/h and wavelength λ

29. Bohr’s Correspondence Principle • Bohr’s Correspondence Principle states that quantum mechanics is in agreement with classical physics when the energy differences between quantized levels are very small • Similar to having Newtonian Mechanics be a special case of relativistic mechanics when v << c

30. Successes of the Bohr Theory • Explained several features of the hydrogen spectrum • Accounts for Balmer and other series • Predicts a value for RH that agrees with the experimental value • Gives an expression for the radius of the atom • Predicts energy levels of hydrogen • Gives a model of what the atom looks like and how it behaves • Can be extended to “hydrogen-like” atoms • Those with one electron • Ze2 needs to be substituted for e2 in equations • Z is the atomic number of the element

31. Modifications of the Bohr Theory – Elliptical Orbits 28.4 • Sommerfeld extended the results to include elliptical orbits • Retained the principle quantum number, n • Determines the energy of the allowed states • Added the orbital quantum number, ℓ • ℓ ranges from 0 to n-1 in integer steps • All states with the same principle quantum number are said to form a shell • The states with given values of n and ℓ are said to form a subshell

32. Modifications of the Bohr Theory – Zeeman Effect • Another modification was needed to account for the Zeeman effect • The Zeeman effect is the splitting of spectral lines in a strong magnetic field • This indicates that the energy of an electron is slightly modified when the atom is immersed in a magnetic field • A new quantum number, mℓ, called the orbital magnetic quantum number, had to be introduced • m ℓ can vary from - ℓ to + ℓ in integer steps

33. Modifications of the Bohr Theory – Fine Structure • High resolution spectrometers show that spectral lines are, in fact, two very closely spaced lines, even in the absence of a magnetic field • This splitting is called fine structure • Another quantum number, ms, called the spin magnetic quantum number, was introduced to explain the fine structure

34. de Broglie Waves (28.5-28.14 OMIT!!!!) • One of Bohr’s postulates was the angular momentum of the electron is quantized, but there was no explanation why the restriction occurred • de Broglie assumed that the electron orbit would be stable only if it contained an integral number of electron wavelengths

35. de Broglie Waves in the Hydrogen Atom • In this example, three complete wavelengths are contained in the circumference of the orbit • In general, the circumference must equal some integer number of wavelengths • 2  r = n λ n = 1, 2, …

36. de Broglie Waves in the Hydrogen Atom, cont • The expression for the de Broglie wavelength can be included in the circumference calculation • me v r = n ħ • This is the same quantization of angular momentum that Bohr imposed in his original theory • This was the first convincing argument that the wave nature of matter was at the heart of the behavior of atomic systems

37. de Broglie Waves, cont. • By applying wave theory to the electrons in an atom, de Broglie was able to explain the appearance of integers in Bohr’s equations as a natural consequence of standing wave patterns • Schrödinger’s wave equation was subsequently applied to atomic systems

38. Quantum Mechanics and the Hydrogen Atom • One of the first great achievements of quantum mechanics was the solution of the wave equation for the hydrogen atom • The significance of quantum mechanics is that the quantum numbers and the restrictions placed on their values arise directly from the mathematics and not from any assumptions made to make the theory agree with experiments

39. Quantum Number Summary • The values of n can range from 1 to  in integer steps • The values of ℓ can range from 0 to n-1 in integer steps • The values of mℓ can range from -ℓ to ℓ in integer steps • Also see Table 28.2

40. Spin Magnetic Quantum Number • It is convenient to think of the electron as spinning on its axis • The electron is not physically spinning • There are two directions for the spin • Spin up, ms = ½ • Spin down, ms = -½ • There is a slight energy difference between the two spins and this accounts for the doublet in some lines

41. Spin Notes • A classical description of electron spin is incorrect • Since the electron cannot be located precisely in space, it cannot be considered to be a spinning solid object • P. A. M. Dirac developed a relativistic quantum theory in which spin appears naturally

42. Electron Clouds • The graph shows the solution to the wave equation for hydrogen in the ground state • The curve peaks at the Bohr radius • The electron is not confined to a particular orbital distance from the nucleus • The probability of finding the electron at the Bohr radius is a maximum

43. Electron Clouds, cont • The wave function for hydrogen in the ground state is symmetric • The electron can be found in a spherical region surrounding the nucleus • The result is interpreted by viewing the electron as a cloud surrounding the nucleus • The densest regions of the cloud represent the highest probability for finding the electron

44. Wolfgang Pauli • 1900 – 1958 • Contributions include • Major review of relativity • Exclusion Principle • Connect between electron spin and statistics • Theories of relativistic quantum electrodynamics • Neutrino hypothesis • Nuclear spin hypothesis

45. The Pauli Exclusion Principle • No two electrons in an atom can ever have the same set of values of the quantum numbers n, ℓ, mℓ, and ms • This explains the electronic structure of complex atoms as a succession of filled energy levels with different quantum numbers

46. Filling Shells • As a general rule, the order that electrons fill an atom’s subshell is: • Once one subshell is filled, the next electron goes into the vacant subshell that is lowest in energy • Otherwise, the electron would radiate energy until it reached the subshell with the lowest energy • A subshell is filled when it holds 2(2ℓ+1) electrons • See table 28.3

47. The Periodic Table • The outermost electrons are primarily responsible for the chemical properties of the atom • Mendeleev arranged the elements according to their atomic masses and chemical similarities • The electronic configuration of the elements explained by quantum numbers and Pauli’s Exclusion Principle explains the configuration

48. Characteristic X-Rays • When a metal target is bombarded by high-energy electrons, x-rays are emitted • The x-ray spectrum typically consists of a broad continuous spectrum and a series of sharp lines • The lines are dependent on the metal of the target • The lines are called characteristic x-rays

49. Explanation of Characteristic X-Rays • The details of atomic structure can be used to explain characteristic x-rays • A bombarding electron collides with an electron in the target metal that is in an inner shell • If there is sufficient energy, the electron is removed from the target atom • The vacancy created by the lost electron is filled by an electron falling to the vacancy from a higher energy level • The transition is accompanied by the emission of a photon whose energy is equal to the difference between the two levels

50. Moseley Plot • λ is the wavelength of the K line • K is the line that is produced by an electron falling from the L shell to the K shell • From this plot, Moseley was able to determine the Z values of other elements and produce a periodic chart in excellent agreement with the known chemical properties of the elements